Lewis structures and molecular shapes

Counting electrons first. Add up the valence electrons contributed by every atom in the molecule or ion; add one electron per unit of negative charge, subtract one per unit of positive charge. This total, divided by two, is the number of electron pairs you have to place — no more, no fewer.

Building the skeleton. The least electronegative non-hydrogen atom is usually placed centrally, with the remaining atoms bonded around it by single bonds. Octets on the terminal atoms are completed first with lone pairs; whatever is left over goes on the central atom.

Fixing an incomplete octet. If the central atom still has fewer than eight electrons after this first pass, a lone pair on a terminal atom is converted into a second (or third) bond pair to the centre, turning a single bond into a double or triple bond. This is exactly how carbon dioxide gets its two C=O double bonds instead of two weak single bonds with an electron-deficient carbon.

Judging the result with formal charge. More than one Lewis structure can often be drawn for the same formula. The formal charge on each atom is the tool used to decide which structure is the more realistic picture of the real bonding:

where: $V$ is the number of valence electrons the free atom would have, $N$ is the number of non-bonding (lone-pair) electrons it owns in the structure, and $B$ is the number of bonding electrons around it (so $\tfrac{1}{2}B$ is its "share" of the bonds). A structure with formal charges close to zero, and any negative charge sitting on the more electronegative atom, is preferred over one with larger or badly placed charges.

Draw the Lewis structure of ozone, $\text{O}_3$, and use formal charge to describe its bonding.

Step 1. Count valence electrons: each O contributes 6, and $\text{O}_3$ is neutral, so total $= 3 \times 6 = 18$ electrons (9 pairs).

Step 2. Build the skeleton with a central O bonded to two terminal O atoms by single bonds: this uses 2 bond pairs (4 electrons), leaving 14 electrons (7 pairs).

Step 3. Complete octets on the two terminal oxygens first: 3 lone pairs (6 electrons) on each, using 12 electrons and leaving only 2 electrons (1 pair) for the central atom. The central O then has 2 bond pairs plus 1 lone pair — only 6 electrons, an incomplete octet.

Step 4. Convert one lone pair from a terminal oxygen into a second bond pair to the centre, turning one O–O single bond into an O=O double bond. Recheck: the central O now has $4+2+2=8$ electrons — a complete octet. The double-bonded terminal O has $4+4=8$; the single-bonded terminal O has $2+6=8$.

Step 5. Compute formal charges with $\text{FC} = V - N - \tfrac{1}{2}B$. Central O: $V=6$, $N=2$, $B=6$, so $\text{FC} = 6 - 2 - 3 = +1$. Double-bonded terminal O: $V=6$, $N=4$, $B=4$, so $\text{FC} = 6 - 4 - 2 = 0$. Single-bonded terminal O: $V=6$, $N=6$, $B=2$, so $\text{FC} = 6 - 6 - 1 = -1$. The three formal charges sum to $+1 + 0 - 1 = 0$, correctly matching a neutral molecule.

Step 6. Since either terminal oxygen could equally well carry the double bond, two equivalent Lewis structures exist, and the true molecule is a resonance hybrid of both. Experimentally, both O–O bonds in ozone are identical, about $1.28\ \text{Å}$ long — intermediate between a typical O–O single bond ($\approx 1.48\ \text{Å}$) and an O=O double bond ($\approx 1.21\ \text{Å}$), exactly as expected for a bond that is "half single, half double" in the hybrid. Answer: ozone needs two resonance structures with formal charges $+1$ (central O), $0$, and $-1$ (terminal O's), and its real bonds are two equal, intermediate-length O–O bonds, not one single and one double.

• Stopping after placing single bonds without checking every non-hydrogen atom for a complete octet — an electron-short central atom almost always needs a lone pair converted to a bond pair.
• Treating resonance as the molecule physically "flipping" between structures over time — the real molecule is one static hybrid, not a fast interconversion.
• Comparing two candidate structures without computing formal charge on every atom — the structure with charges closest to zero (and negative charge on the more electronegative atom) is the better picture, not just "whichever looks neater."
• Letting resonance structures move an atom's position — only electrons move between resonance forms; the nuclear skeleton stays fixed.

Ask the tutor: Want to work through the Lewis structure and formal charges of a molecule like $\text{SO}_2$ or $\text{NO}_3^-$ together, one electron pair at a time?